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Determining Molarity Through Acid-Base Titration - Lab Report
General Chemistry Lab (CHEM 1203L)
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Determining Molarity Through Acid-Base Titration
Joshua Farley CHEM 1251L- 10/30/ Introduction This experiment focused on an essential quantitative technique that, when used effectively, can determine the concentration of an acid in a solution. This process is known as titration, or volumetric analysis. The purpose of this particular lab experiment was to effectively use titration by initially standardizing the base sodium hydroxide, NaOH, by determining its exact molarity so that it could act as the titrant throughout the rest of the experiment. The titrant is a solution of known concentration that, in this experiment, was funneled into a buret in order to precisely and slowly find the exact equivalence point—the point in which all of the analyte had been used up and therefore, both reactants became the limiting reagents. However, the experimental end of the titration, known as the end point, is signaled by a very weak acid or base called an indicator. The indicator chosen for this experiment is phenolphthalein, a weak, colorless acid that releases an acidic proton when the neutralization has been completed. This proton will leave the bright pink anion (In-) free of its attractions, in which the solution will also turn a pink color for chemists to observe.
So why can we not just weigh out the NaOH in solid form and convert the value to moles? Well, this base is hygroscopic, meaning that it immediately absorbs all surrounding moisture from the air. So by the time you extract the sodium hydroxide and get it to the scale, it
will already be heavily contaminated with water vapor in the air. This makes NaOH a secondary standard, which basically means that its concentration could only be determined by comparing it to that of a primary standard, which is a substance of high purity and stability. In this experiment, the primary standard was an acid called potassium hydrogen phthalate, KHC 8 H 4 O 4 , more easily remembered as KHP. The chemical formula that represented this acid-base neutralization reaction is: KHC 8 H 4 O 4 (aq) + NaOH (aq) KNaC 8 H 4 O 4 (aq) + H 2 O (l). Notice that the molar ratio between KHP and sodium hydroxide is 1:1. This means that the number of moles of KHP will be the same as the number of moles of sodium hydroxide because 1 mole of KHP reacts with 1 mole of NaOH to carry out the reaction, leaving no excess reagents. It is important to pay close attention to the molar ratios because although the monoprotic acid (acetic acid) also maintains a 1:1 mole-to-mole ratio with NaOH, a diprotic acid, such as sulfuric acid, has a 1: acid-to-base ratio. In this case, 2 moles of NaOH are required to titrate 1 mole of H 2 SO 4.
Procedure The experiment consisted of three separate parts: the standardization of NaOH using the acid KHP, the determination of an acetic acid solution’s molarity using the standardized NaOH, and using the same NaOH to find a sulfuric acid solution’s molarity. Part one was critical because without finding an average molarity of NaOH, further calculations would not be possible for part two and three. As stated in the introduction, this step must be done because sodium hydroxide is a secondary standard with hygroscopic properties that make its mass nearly impossible to determine directly with a scale.
The first step in finding the molarity of the sodium hydroxide solution was to weigh about approximately 1 gram of KHP and carefully dissolve it in an Erlenmeyer flask with about 70 mL of distilled water. After recording the measured mass of the substance, the number of
enough NaOH. The exact same titration procedure applied in the same way for all 3 sections of this experiment. After determining the volume of NaOH required to titrate the acetic acid solution, further calculations and observations revealed the molarity of unknown acetic acid ID #138 to be about 1 moles per liter. Because the molar ratio was 1:1 between the acid and the base for this reaction, the same numbers of moles of sodium hydroxide were used to titrate an equal number of moles of acetic acid. This is due to the fact that acetic acid—HC 2 H 3 O 2 —is a monoprotic acid and only has one ionizable H+ ion per molecule, which relates to NaOH, which only has one ionizable OH- ion. This part of the experiment was repeated twice to ensure accuracy once again and recognize any possible errors that could affect the lab.
The third and final part of this titration experiment dealt with a diprotic acid, sulfuric acid. This titration was no different in procedure from the others except that it had an acid-to- base molar ratio of 1:2, meaning that it took 2 moles of NaOH to react with 1 mole of H 2 SO 4. The phenolphthalein was dropped into the sulfuric acid solution 5 times and the process was repeated. As before, distilled water was sprayed on the sides to contain all of the sulfuric acid in the solution while base was being gradually poured in through the buret. After observing the end point of the reaction, the volume of base used was determined and compared to the molarity of NaOH to find the number of moles of NaOH used. The 1:2 mole-to-mole ratio discussed earlier was taken into perspective to find how many moles of sulfuric acid was reacted. The average molarity for this value came out to be about 0 moles per liter of solution.
Results
To find the molar mass of potassium hydrogen phthalate—KHC 8 H 4 O 4 —the atomic mass of each element was added together.
K=39 g/mol H=1 g/mol O=16 g/mol 39 g/mol + 5(1 g/mol) + 8(12 g/mol) + 4(16 g/mol) = 204 g/mol
Table 1: All essential data recorded and calculated during part 1, the standardization of NaOH.
Data Trial 1 Trial 2 Trial 3 Mass KHP Used (g) 1 g 1 g 1 g Molar Mass of KHP 204 g/mol 204 g/mol 204 g/mol Moles of KHP Used 0 mol 0 mol 0 mol Acid:Base Molar Ratio 1:1 1:1 1: Moles of NaOH 0 mol 0 mol 0 mol Initial Volume of NaOH 2 mL 7 mL 2 mL Final Volume of NaOH 27 mL 34 mL 26 mL Volume of NaOH Used 24 mL, 0. L
27 mL, 0. L
24 mL, 0. L Molarity of NaOH 0 M 0 M 0 M Average Molarity of NaOH
0 M
Equation 1 shows how to find the volume used in a solution while operating the buret. Vf represents the volume recorded at the end (final volume) while Vi represents the initial volume.
Eq. 1) Volume = Vf - Vi
Discussion In this titration experiment, a very crucial aspect to take into account were the chemical formulas for each reaction because they illustrate exactly what will occur and it showed the acid- base molar ratio. The chemical formulas critical to this lab are as follow:
- KHC 8 H 4 O 4 (aq) + NaOH (aq) KNaC 8 H 4 O 4 (aq) + H 2 O (l)
- HC 2 H 3 O 2 (aq) +NaOH (aq) NaC 2 H 3 O 2 (aq) + H 2 O (l)
- H 2 SO 4 (aq) + 2NaOH (aq) Na 2 SO 4 (aq) + H 2 O (l) Reaction 1 illustrates the titration reaction between a known amount of titrant (KHP) and an unknown amount of analyte (NaOH). This reaction was used to standardize NaOH and determine its exact molarity. This had to be done right because the molarity of sodium hydroxide found was to be used throughout the remainder of the experiment. Without the concentration of sodium hydroxide, the number of moles of base would be impossible to find, which would also prevent the moles of the acid, either monoprotic or diprotic, from being found. This standardization was repeated 3 times in order to ensure that the molarity would be determined as accurately as possible. Despite this, there were still several minute errors that may have directly affected our results. One error that may have easily occurred in several of the trials had to do with the amount of base being added to find an end point. The observance of the endpoint always occurred so quickly that it was often difficult to ensure that the exact amount of base needed was added. Additional drops of sodium hydroxide could have been added in excess and ultimately changed the outcome ever so slightly. When the end point has been reached, the entire solution instantly converts to a light pink color and any additional base will continue to darken the shade of pink. Several of the trials ended with a solution that was a slightly darker shade of pink than expected.
This was not the only possible experimental error observed. During the actual act of performing a titration, it is important to agitate the solution (typically by stirring it) and also to rinse the walls of the Erlenmeyer flask with distilled water occasionally. This is essential because some of the titrant used can get stuck to the sides and this “loss” of titrant will result in a lower actual molarity than what it should be.
The equivalence point is the exact point at which the moles of titrant is stoichiometrically equal to the moles of analyte; also known as the theoretical end of a titration. Unfortunately, this point could not be directly observed with the human eye and a signal, or indicator, was used to find the end point. The end point is similar to the equivalence point except that it is the experimental end of a titration reaction. The end point is signaled by using an indicator, such as phenolphthalein. This indicator releases its anions at a specific pH which will result in a pink solution once it has neutralized completely. Since some salts can form that will affect the pH, the indicator changes in a pH range of 8.0-9. Although the equivalence point and end point are very close together, they still have a little bit of distance from one another and, therefore, the recorded volumes were not 100% accurate to the equivalence points.
Reaction 2 represents the titration between acetic acid, which is a monoprotic acid, and the titrant, sodium hydroxide. Reaction 3 shows a titration reaction between sulfuric acid, a diprotic acid, and the same titrant, sodium hydroxide. The key difference between these two reactions is that acetic acid is monoprotic, meaning it only has 1 ionizable H+ ion. Sulfuric acid, on the other hand, has 2 ionizable H+ ions, making it diprotic. So what does this mean? Well, once the equations have been balanced out, one can see that acetic acid only requires 1 mole of NaOH to titrate it. This results in a simple 1:1 molar ratio. However, sulfuric acid requires 2 moles of NaOH to titrate it; a 1:2 molar ratio. To make it simple, if the concentrations of acetic
Based on many of the values shown in the tables shown above, it appears that the experiment maintained mostly accurate data. One way to determine this accuracy is to look at the moles of NaOH used per trial in comparison with the volume of NaOH used in milliliters. This values remain proportional to the molarity. So, although the experiment showed several signs of significant error, the outcome stayed relatively accurate and precise with all other data collected.
Sample Calculations
Finding the molar mass of Potassium hydrogen phthalate. K=39 g/mol H=1 g/mol O=16 g/mol 39 g/mol + 5(1 g/mol) + 8(12 g/mol) + 4(16 g/mol) = 204 g/mol
Finding moles of KHP used: molarmassgrams 1 g / 204 g/mol = 0 moles
Finding volume of NaOH used: V=Vf−Vi V = 27 mL – 2 mL V = 24 mL
Finding the molarity of any substance: M=litersof solutionmolesof solute
M = 0 moles / 0 L
M = 0 moles/L
5) Finding the average molarity between 2+ values: ∑¿of molarityvaluesof valuesadded
(0 M + 0 M + 0 M) / 3
= 0 M
Determining Molarity Through Acid-Base Titration - Lab Report
Course: General Chemistry Lab (CHEM 1203L)
University: University of North Carolina at Charlotte
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